1. 1 Vanessa N. Prasad-Permaul Valencia Community College CHM 1045

2. 2 Ionic Bonds  Ionic Bonds: a chemical bond formed by the electrostatic attraction between positive and negative ions  Bond forms when one or more electrons are transferred from the valence shell of one atom to the valence shell of the other.  Cation: atom losing electrons  Anion: atom gaining electrons 2

3. 3 Ionic Bonds Na + Cl Na + + Cl - [Ne]3s 1 + [Ne]3s 2 3p 5 [Ne] + [Ne]3s 2 3p 6 3

4.  Lewis Electron-Dot Symbols: electrons in the valence shell of an atom or ion are represented by dots placed around the elemental symbol. Na + Cl Na + + [ Cl ] -  Valence electrons are those electrons with the highest principal quantum number (n). 4 Ionic Bonds

5. 5 Lewis Electron-Dot Symbols for Atoms in the Periodic Table of Elements

6. 6 Ionic Bonds EXAMPLE 9.1: Use Lewis Electron-Dot symbols to represent the transfer of electrons form magnesium to fluorine atoms to form ions with noble gas configuration. Mg + Fl MgFl 2 F + Mg + F [ F ] - + Mg 2+ + [ F ] -

7. 7 Ionic Bonds EXERCISE 9.1: Use Lewis Electron-Dot symbols to represent the transfer of electrons form magnesium to oxygen atoms to form ions with noble gas configuration.

8. 8 Electron Configurations of Ions Common monatomic ions found in compounds of the main-group elements fall into 3 categories: 1.Cations of Groups IA to IIIA having noble gas or pseudo-noble-gas configurations ; the ion charges equal group numbers. 2.Cations of Groups IIIA to VA having the ns 2 electrons configurations; the ion charges equal the group numbers minus two. Tl +, Sn 2+, Pb 2+, Bi 3+. 3.Anions of Groups VA to VIIA having noble gas configurations; the ion charges equal the group number minus 8.

9. 9 EXAMPLE 9.2: Write the electron configuration and Lewis symbol for N 3-. N = [He]2s 2 2p 3 + 3e - [He]2s 2 2p 6 N = N N 3- = [ N ] 3- Electron Configurations of Ions

10. 10 EXERCISE 9.2: Write the electron configuration and Lewis symbol for Ca 2+ and S 2-. Electron Configurations of Ions

11. 11 EXERCISE 9.3: Write the electron configuration and Lewis symbol for Pb and Pb 2+. Electron Configurations of Ions

12. 12 Electron Configurations of Ions EXAMPLE 9.3: Write the electron configuration and Lewis symbol for Fe 2+ and Fe 3+. Fe = [Ar]3d 6 4s 2 = Fe Fe 2+ = [Ar]3d 6 = [ Fe ] 2+ Fe 3+ = [Ar]3d 5 = [ Fe ] 3+

13. 13 Electron Configurations of Ions EXERCISE 9.4: Write the electron configuration and Lewis symbol for Mn and Mn 2+.

14. 14 Ionic Radii Ionic Radius: A measure of the size of the spherical region around the nucleus of an ion within the electrons are most likely to be found.  If the atom loses an electron, the cation will be smaller.  The electron-electron repulsion is initially less.  orbitals can shrink to increase the attraction of the electrons for the nucleus.

15. 15 Ionic Radii EXERCISE 9.5: Which has the larger radius. S or S 2-. Explain.

16. 16 Ionic Radii EXERCISE 9.6: Using only the Periodic Table, arrange the following ions in order of increasing ionic radius : Sr 2, Mg 2+, Ca 2+.

17. 17 Ionic Radii The Comparison of Atomic and Ionic Radii DECREASES ACROSS A PERIOD INCREASESDOWN AGROUPINCREASESDOWN AGROUP ATOMIC RADIUS

18. 18 Ionic Radii Isoelectronic: Different species having the same number and configuration of electrons. Na + < Mg 2+ < Al 3+ Decrease in atomic radius  As the charge increases, the orbitals contract due to the greater attractive forces of the nucleus  the ionic radius decreases with increasing atomic number.  In general, across a period the cations decrease in radius.  As the anions are reached, there is an abrupt increase in radius and then the radius decreases again.

19. 19 Ionic Radii EXAMPLE 9.4: Arrange the following ions in order of decreasing ionic radius: F -, Mg 2+, O 2-. F - = 1s 2 2s 2 2p 6 Mg 2+ = 1s 2 2s 2 2p 6 O 2- = 1s 2 2s 2 2p 6 All are isoelectronic, so as the nuclear charge increases, the ionic radius decreases. O 2-, F -, Mg 2+

20. 20 Ionic Radii EXERCISE 9.7: Arrange the following ions in order of increasing ionic radius: Cl -, Ca 2+, P 3-.

21. Covalent Bonds  Covalent Bonds are formed by sharing at least one pair of electrons.  The attraction (nucleus/electrons) outweighs the repulsions (electron/electron & nucleus/nucleus) 21

22. Covalent Bonds 22 Every covalent bond has a characteristic length that leads to maximum stability: Bond Length

23. Strength of Covalent Bonds 23 Energy required to break a covalent bond in an isolated gaseous molecule is called the bond dissociation energy. Same amount of energy released when the bond forms.

24. Electron-Dot Structures  The electron-dot structures provide a simple, but useful, way of representing chemical reactions.  Ionic:  Covalent: 24

25. 25 Covalent Bonds Coordinate covalent bond: a bond formed when both electrons of the bond are donated by one atom A + B A B H + H + + NH3 H N H H

26. 26  Group 1A tends to lose their ns 1 valence shell electron to adopt a noble gas electron configuration.  Group 2A lose both ns 2  Group 3A lose all three ns 2 np 1  Group 7A Gains one electron to attain noble gas configuration  Group 8A inert, rarely lose or gain electrons Covalent Bonds OCTET RULE

27. Single Bonds: Double Bonds: Triple Bonds: 27 Covalent Bonds

28. 28 Covalent Bonds

29. 29 Covalent Bonds ELECTRONEGATIVITY: a measure of the ability of an atom in a molecule to draw bonding electrons to itself.  Bond polarity is due to electronegativity differences between atoms.  Pauling Electronegativity: is expressed on a scale where F = 4.0

30. Pauling Electronegativities 30 Covalent Bonds

31. 31 Covalent Bonds EXAMPLE 9.5: Use electronegativity values to arrange the following bonds in order of increasing polarity: P H, H O, C Cl. P H = 0.0 H O = 1.4 C Cl = 0.5 P H, C Cl, H O

32. 32 Covalent Bonds EXERCISE 9.8: Use electronegativity values to arrange the following bonds in order of increasing polarity: C O, C S, H Br.

33. Drawing Lewis-Dot Structures Rule 1: Count the total valence electrons. Rule 2: Draw the structure using single bonds. Rule 3: Distribute the remaining electron pairs around the peripheral atoms. Rule 4: Put remaining pairs on central atom. Rule 5: Share lone pairs between bonded atoms to create multiple bonds. 33

34. Drawing Lewis-Dot Structures 34

35. Lewis-Dot Structures  NH 2 F Amino Fluoride: In this molecule, nitrogen is the central atom.  Rule 1: Number of electrons = 5 + (2 x 1) + 7 = 14 = 7 pairs 35 Rule 2Rule 3Rule 4

36. 36 EXAMPLE 9.6: Sulfur dichloride SCl 2, is a red fuming liquid used in the manufacture insecticides. Write the Lewis formula for the molecule. S = 6 Cl = 7 each for a total of 20 electrons :Cl : S : Cl: :Cl S Cl: Lewis-Dot Structures ¨¨ ¨¨ ¨ ¨ ¨¨ ¨¨ ¨ ¨

37. 37 EXERCISE 9.9: Dichlorodifluoromethane CCl 2 F 2, is a gas used as a refrigerant and aerosol propellant. Write the Lewis formula for the molecule. Lewis-Dot Structures

38. 38 EXAMPLE 9.7: Carbonyl chloride or phosgene, COCl 2, is a highly toxic gas used as a starting material for the preparation of polyurethane plastics. What is the electron dot structure of this compound? C = 4, O = 6, Cl =7 each for a total of 24 electrons : Cl : C : Cl : :O: : Cl C Cl : :O: Lewis-Dot Structures ¨¨ ¨¨ ¨ ¨ ¨ ¨ ¨ ¨

39. 39 EXERCISE 9.10: Write the electron-dot structure of carbon dioxide. Lewis-Dot Structures

40. 40 EXAMPLE 9.8: Obtain the electron-dot formula for the BF 4 - ion. B = 3, F = 7 each (7 x 4) = 28 for a total 31 electrons. It is an ion with one more electron so a total of 32 electrons. : F : : F : B : F : : F : Lewis-Dot Structures ¨ ¨ ¨¨ ¨ ¨ ¨ ¨ : F : : F B F : : F : ¨ ¨ ¨¨ ¨ ¨

41. 41 EXERCISE 9.11: Write the electron-dot structure of: A. the hydronium ion, H 3 O + B. The chlorite ion, ClO 2 - Lewis-Dot Structures

42. Resonance Structures  When multiple structures can be drawn, the actual structure is an average of all possibilities.  The average is called a resonance hybrid. A straight double-headed arrow indicates resonance. 42 OOO OOO

43. 43 EXAMPLE 9.9: Describe the electron structure of the carbonate ion CO 3 2-, in terms of electron-dot formulas. C = 4, O = 3 x 6 = 18 for a total of 22 electrons, but it has gained two electrons so there is a total of 24 electrons. 2- 2- 2- : O : : O : : O : C C C : O : : O : : O : O : : O : O : Lewis-Dot Structures ¨¨¨¨¨¨ ¨¨

44. 44 EXERCISE 9.12: Describe the bonding in NO 3 - using resonance formulas. Lewis-Dot Structures

45. Formal Charge  Formal Charge: Determines the best resonance structure.  We determine formal charge and estimate the more accurate representation. Formal Charge = valence e - - # of e - in a bond - (# of lone-pair e - ) 2 45

46. 46 ¨ ¨ ¨ ¨ Formal Charge :Cl C Cl: :O: ¨ ¨ ¨ ¨ :Cl C Cl :O: ¨ ¨ ¨ ¨ Cl C Cl: :O: Cl = 7 – (2/2) – 6 = 0 O = 6 – (4/2) – 4 = 0 C = 4 – (8/2) – 0 = 0 Cl = 7 – (4/2) – 4 = +1 O = 6 – (2/2) – 6 = -1 C = 4 – (8/2) – 0 = 0 ¨ ¨ + - - + Cl = 7 – (4/2) – 4 = +1 O = 6 – (2/2) – 6 = -1 C = 4 – (8/2) – 0 = 0

47. 47 EXAMPLE 9.11: Write the Lewis formula that best describes the charge distribution in the sulfuric acid molecule, H 2 SO 4, according to the rules of formal charge. :O: H O S O HH :O: Formal Charge ¨ ¨ ¨ ¨ ¨ ¨ +2 ¨ O ¨ SOH ¨ ¨ O O ¨ ¨ ¨ ¨

48. 48 Exercise 9.15:Write the Lewis formula that best describes the phosphoric acid molecule, H 3 PO 4.

49. Resonance Structures  How is the double bond formed in O 3 ?  The correct answer is that both are correct, but neither is correct by itself. 49

50. Example 1: Which of the following is correct? 1. Energy is absorbed to form a bond 2. Energy is released when a bond is formed 50

51. Example 2: Drawing Lewis-Dot Structures  Draw electron-dot structures for: C 3 H 8 H 2 O 2 CO 2 N 2 H 4 CH 5 NC 2 H 4 C 2 H 2 Cl 2 CO H 3 S + HCO 3 – 51

52. Example 3: Formal Charge  Calculate the formal charge and determine the most favorable of the following electron dot structures: SO 2 NO 3 – NCO – N 2 OO 3 CO 3 2– 52

53. Example 4: What is the overall formal charge of the following structure? 1. -2 2. -3 3. -1 4. 0 53

54. Example 5: Ionic Radii of Ions  Compare ionic radii  Fe & Fe 3+  Cl & Cl - 54